LINUS PAULING (1901- 94)

1931 USA

‘A framework for understanding the electronic and geometric structure of molecules and crystals’

An important aspect of this framework is the concept of hybridisation: in order to create stronger bonds, atoms change the shape of their orbitals (the space around a nucleus in which an electron is most likely to be found) into petal shapes, which allow more effective overlapping of orbitals.

A chemical bond is a strong force of attraction linking atoms in a molecule or crystal. BOHR had already shown that electrons inhabit fixed orbits around the nucleus of the atom. Atoms strive to have a full outer shell (allowed orbit), which gives a stable structure. They may share, give away or receive extra electrons to achieve stability. The way atoms will form bonds with others, and the ease with which they will do it, is determined by the configuration of electrons.

Earlier in the century, Gilbert Lewis (1875-1946) had offered many of the basic explanations for the structural bonding between elements, including the sharing of a pair of electrons between atoms and the tendency of elements to combine with others to fill their electron shells according to rigidly defined orbits (with two electrons in the closest orbit to the nucleus, eight in the second orbit, eight in the third and so on).

Pauling was the first to enunciate an understanding of a physical interpretation of the bonds between molecules from a chemical perspective, and of the nature of crystals.

In a covalent bond, one or more electrons are shared between two atoms. So two hydrogen atoms form the hydrogen molecule, H2, by each sharing their single electron. The two atoms are bound together by the shared electrons. This was proposed by Lewis and Irving Langmuir in 1916.

In an ionic bond, one atom gives away one or more electrons to another atom. So in common salt, sodium chloride, sodium gives away its spare electron to chlorine. As the electron is not shared, the sodium and chlorine atoms are not bound together in a molecule. However, by losing an electron, sodium acquires a positive charge and chlorine, by gaining an electron, acquires a negative charge. The resulting sodium and chlorine ions are held in a crystalline structure. Until Pauling’s explanation it was thought that they were held in place only by electrical charges, the negative and positive ions being drawn to each other.

Pauling’s work provided a value for the energy involved in the small, weak hydrogen bond.
When a hydrogen atom forms a bond with an atom which strongly attracts its single electron, little negative  charge is left on the opposite side of the hydrogen atom. As there are no other electrons orbiting the hydrogen nucleus, the other side of the atom has a noticeable positive charge – from the proton in the nucleus. This attracts nearby atoms with a negative charge. The attraction – the hydrogen bond – is about a tenth of the strength of a covalent bond. 
In water, attraction between the hydrogen atoms in one water molecule and the oxygen atoms in other water molecules makes water molecules ‘sticky’. It gives ice a regular crystalline structure it would not have otherwise. It makes water liquid at room temperature, when other compounds with similarly small molecules are gases at room temperature.Water10_animation

One aspect of the revolution he brought to chemistry was to insist on considering structures in terms of their three-dimensional space. Pauling showed that the shape of a protein is a long chain twisted into a helix or spiral. The structure is held in shape by hydrogen bonds.
He also explained the beta-sheet, a pleated sheet arrangement given strength by a line of hydrogen bonds.

He devised the electronegativity scale, which ranks elements in order of their electronegativity – a measure of the attraction an atom has for the electrons involved in bonding (0.7 for caesium and francium to 4.0 for fluorine). The electronegativity scale lets us say how covalent or ionic a bond is.

Pauling’s application of quantum theory to structural chemistry helped to establish the subject. He took from quantum mechanics the idea of an electron having both wave-like and particle-like properties and applied it to hydrogen bonds. Instead of there being just an electrical attraction between water molecules, Pauling suggested that wave properties of the particles involved in hydrogen bonding and those involved in covalent bonding overlap. This gives the hydrogen bonds some properties of covalent bonds.

1922 – while investigating why atoms in metals arrange themselves into regular patterns, Pauling used X-ray diffraction at CalTech to determine the structure of molybdenum.

When X-rays are directed at a crystal, some are knocked off course by striking atoms, while others pass straight through as if there are no atoms in their path. The result is a diffraction pattern – a pattern of dark and light lines that reveal the positions of the atoms in the crystal.
Pauling used X-ray and electron diffraction, magnetic effects and measurements of the heat of chemical reactions to calculate the distances and angles between atoms forming bonds. In 1928 he published his findings as a set of rules for working out probable crystalline structures from the X-ray diffraction patterns.

1939 – ‘The Nature of the Chemical Bond and the Structure of Molecules’
Pauling suggests that in order to create stronger bonds, atoms change the shapes of their waves into petal shapes; this was the ‘hydridisation of orbitals’.
Describing hybridisation, he showed that the labels ‘ionic’ and ‘covalent’ are little more than a convenience to group bonds that really lie on a continuous spectrum from wholly ionic to wholly co-valent.

Pauling developed six key rules to explain and predict chemical structure. Three of them are mathematical rules relating to the way electrons behave within bonds, and three relate to the orientation of the orbitals in which the electrons move and the relative position of the atomic nuclei.

          

   

1951 – published his findings one year after WILLIAM LAWRENCE BRAGG’s team at the Cavendish Laboratory.

CARBON BONDING
As carbon has four filled and four unfilled electron shells it can form bonds in many different ways, making possible the myriad organic compounds found in plants and animals. The concept of hybridisation proved useful in explaining the way carbon bonds often fall between recognised states, which opened the door to the realm of organic chemistry.

X-ray diffraction alone is not very useful for determining the structure of complex organic molecules, but it can show the general shape of the molecule. Pauling’s work showed that physical chemistry at the molecular level could be used to solve problems in biology and medicine.

  

A problem that needed resolving was the distance between particular atoms when they joined together. Carbon has four bonds, for instance, while oxygen can form two.It would seem that in a molecule of carbon dioxide, which is made of one carbon and two oxygen atoms, two of carbon’s bonds will be devoted to each oxygen.

diagram of CO2 bond length

CO2 bond length

Well-established calculations gave the distance between the carbon and oxygen atoms as 1.22 × 10-10m. Analysis gave the size of the bond as 1.16 Angstroms. The bond is stronger, and hence shorter. Pauling’s quantum .3-2. explanation was that the bonds within carbon dioxide are constantly resonating between two alternatives. In one position, carbon makes three bonds with one of the oxygen molecules and has only one bond with the other, and then the situation is reversed.

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SIR JAMES CHADWICK (1891-1974)

1932 Manchester, England

‘Discovery of neutrons – elementary particles devoid of any electric charge’

In contrast with the Helium nuclei (alpha rays) which are charged, and therefore repelled by the electrical forces present in the nuclei of heavy atoms, the neutron is capable of penetrating and splitting the nuclei of even the heaviest elements, creating the possibility of the fission of 235uranium

Assistant to ERNEST RUTHERFORD, Chadwick’s earlier work involved the showering of elements with alpha particles. The picture that gradually emerged was one of a nucleus that contained a very heavy particle with a positive electric charge. This particle was christened the proton, the hydrogen building block envisaged by WILLIAM PROUT.
A spin-off of this was the deduction that the nucleus of the hydrogen atom, the positively charged proton with an atomic weight of one, was present in larger quantities in the nucleus of every other atom.

Rutherford and Geiger had shown that a helium atom and an alpha particle were the same thing, apart from the positive electric charge carried by the alpha particle.

A helium atom seemed to consist of a nucleus of a pair of protons circled by two electrons. However, a helium nucleus seemed to weigh as much as four protons. The mass of the known components of an atom did not add-up. Protons seemed to account for around half of the weight and were matched in number by an equal amount of negatively charged electrons to counter their positive charge. But the weight of an electron was one-thousandth that of a proton, so approximately half of the atomic weight of the element was unaccounted for.
Chadwick solved the conundrum in 1932 when he re-interpreted the results of an experiment carried out by IRENE and FREDERIC JULIOT-CURIE (Irene was the daughter of PIERRE and MARIE CURIE).
The couple had found in 1932 that when beryllium was showered with alpha particles, the resultant radiation could force protons out of substances containing hydrogen. Chadwick suggested that neutrally charged sub-atomic units, which he named neutrons, with the same weight as protons, could force this reaction and therefore were what made up the radiation that the Curies called gamma rays. Rutherford had hinted at the existence of such a particle in 1920.

The explanation was widely accepted and the riddle of `atomic weight’ had been solved: a similar number of neutrons to protons in the nucleus of an element would make up the remaining fifty per cent of the previously ‘missing’ mass.

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FREDERICK SODDY (more)

The discovery of the neutron made sense of the observation that many elements come in a variety of forms, each with differing radioactive properties such as decay rate. Each form consisted of atoms with a different mass. Frederick Soddy christened these variants ‘isotopes’ in 1911. The idea that each element might be a mixture of atoms of different atomic weights explained why the atomic weights of a handful of elements were not simple multiples of the atomic weight of hydrogen, the most notorious example being chlorine whose atomic weight was 35.5 times that of hydrogen. Most of the variant forms of each element turned out to be radioactively unstable. An element such as chlorine, with more than one stable isotope, is rare.

The various isotopes of an element were merely atoms with the same number of protons in their nucleus but with a different number of neutrons.

artistic representation of atomic disintegration

Thus every atom was composed of electrons, protons and neutrons. The protons and neutrons clung together in a central clump – the atomic nucleus – while the electrons circled in a distant haze. The neutrons were responsible for increasing the weight of the elements without adding any electrical charge. Two protons and two neutrons made a helium nucleus; eight protons and eight neutrons an oxygen nucleus; 26 protons and 30 neutrons an iron nucleus; 79 protons and 118 neutrons a gold; and 92 protons and 146 neutrons a nucleus of uranium. When a radioactive nucleus expelled an alpha particle, it lost two neutrons and two protons and consequently became a nucleus of an element two places lower in the periodic table. When a radioactive nucleus emitted a beta particle, however, a neutron changed into a proton, transforming the nucleus into that of an element one place higher in the periodic table.

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FRANCIS CRICK (USA 1916-2004) JAMES DEWEY WATSON (UK b.1928)

1953 – UK

‘The self reproducing genetic molecule DNA has the form of a double helix’

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WATSON & CRICK

The structure explains how DNA stores information and replicates itself.
The helical strands of DNA (deoxyribonucleic acid) consist of chains of alternating sugar and phosphate groups. Four types of base – adenine (A), cytosine (C), guanine (G) and thymine (T) – form the rungs of the DNA ladder, which can only be linked by hydrogen bonds in four combinations: A-T, C-G, T-A, G-C.

The DNA code is based on the order of these four bases and is carried from one generation to the next. The sequence of base pairs along the length of the strands is not the same in DNAs of different organisms. It is this difference in the sequence that makes one gene different from another.

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DMITRI MENDELEEV (1834-1907)

1869 – Russia

‘The properties of elements are periodic functions of their atomic weights’

Arrange the atoms in order of their atomic weight (relative atomic mass) and elements are also arranged in order of their properties. This arrangement of the elements is called the periodic table.

In the modern periodic table elements are no longer arranged by their atomic weight but by a more fundamental quantity; ‘atomic number’.

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DIMITRY IVANOVICH MENDELEYEV

The atomic number of an element is the number of protons in the nucleus of one of its atoms; the number of neutrons, which contributes to atomic weight, is ignored. The modern periodic law is that ‘The properties of elements are periodic functions of their atomic numbers’.

In 1860 Dmitri Ivanovich Mendeleev attended a chemistry conference in Karlsruhe where the Italian Stanislao Cannizzaro’s speech announcing his rediscovery of the distinction between atoms and molecules (originally announced in 1811 by AVOGADRO) made a profound impression.

The German chemist Johann Wolfgang Döbereiner (1780-1849) had recognised mathematical patterns in elements that had similar properties. He found that adding the atomic weights of calcium (40) and barium (137) and dividing the total in two left a value close to the weight of strontium (88). Finding this same pattern repeated for lithium, sodium and potassium, and for chlorine, bromine and iodine confirmed the relationship, which he termed the Law of Triads.

In 1862, French scientist Alexandre Beguyer de Chancourtois developed a way of representing the elements by wrapping a helical list around a cylinder.

A repeating pattern in natural phenomena is a strong indication that there exists a simple, compact description.
The periodic table suggests that the distinct atoms of the elements may be described in terms of significantly fewer building blocks than the number of the individual elements. Atoms, then, were made of significantly fewer subatomic building blocks.

In 1869 the 35-year-old Mendeleev published a table of the 61 elements then known. His list of elements – ‘On the Relation of the Properties to the Atomic Weights of Elements’ – occupied a grid where the atomic weight increased as you went down a column (periods) and the elements in any particular row (groups or families) shared similar properties and valencies (metals and gases, for instance).

Mendeleev had to juggle the order of a few elements, assuming their weights to have been incorrectly measured, and predicted that some undiscovered elements would fill the gaps in the table, based on the properties of the elements surrounding the gaps.
The modern periodic table has been turned sideways.

By 1886, with the discoveries of gallium, scandium and germanium with the properties he had foretold, his prediction was fulfilled. By 1925, chemists had successfully identified all the 92 elements they believed to exist in nature. The first artificial element, neptunium, was synthesised in 1940. Many more elements have been made since then.

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ALEXANDER FLEMING (1881-1955)

1929 – UK

‘First identification of an antibiotic – the discovery of penicillin’

The chance discovery of a mould in 1928 led to the development of a non-toxic drug, which is used to combat the bacteria that infect wounds.

Whilst Paul Erlich (1854-1915) worked in Germany to produce a ‘magic-bullet’, a compound or dye that could stick to bacteria and damage them, Alexander Fleming’s chance discovery of the antibacterial properties of the mould Penicillium notatum led him to conclude there was a chemical produced by the mould that would attack the bacterial agents of disease.

Whilst searching for a naturally occurring bacteria-killer, Fleming’s experiments were concentrated on the body’s own sources, tears, saliva and nasal mucus.
The chance discovery of the anti-bacterial properties of Penicillium notatum was not developed commercially until World War Two over a decade later.

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JOHN DALTON

1808 – England

‘All matter is made up of atoms, which cannot be created, destroyed or divided. Atoms of one element are identical but different from those of other elements. All chemical change is the result of combination or separation of atoms’

Dalton struggled to accept the theory of GAY-LUSSAC because he believed, as a base case, that gases would seek to combine in a one atom to one atom ratio (hence believing the formula of water to be HO not H2O). Anything else would contradict Dalton’s theory on the indivisibility of the atom, which he was not prepared to accept.

The reason for the confusion was that at the time the idea of the molecule was not understood.
Dalton believed that in nature all elementary gases consisted of indivisible atoms, which is true for example of the inert gases. The other gases, however, exist in their simplest form in combinations of atoms called molecules. In the case of hydrogen and oxygen, for example, their molecules are made up of two atoms, described as H2 and O2 respectively.

Gay-Lussac examined various substances in which two elements form more than one type of compound and concluded that if two elements A and B combine to form more than one compound, the different masses of A that combine with a fixed mass of B are in a simple whole number ratio. This is the law of multiple proportions.

AVOGADRO’s comprehension of molecules helped to reconcile Gay-Lussac’s ratios with Dalton’s theories on the atom.

Gay-Lussac’s ratio for water could be explained by two molecules of hydrogen (four ‘atoms’) combining with one molecule of oxygen (two ‘atoms’) to result in two molecules of water (2H2O).

2H2 + O2 ↔ 2H2O

When Dalton had considered water, he could not understand how one atom of hydrogen could divide itself (thereby undermining his indivisibility of the atom theory) to form two particles of water. The answer proposed by Avogadro was that oxygen existed in molecules of two and therefore the atom did not divide itself at all.

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